TRT 401 PHYSICAL CHEMISTRY
SECOND LAW
Second Law Thermodynamic
1.
2.
3.
4.
The direction of Spontaneous Change:
The dispersal of energy
Entropy
Impact on engineering: refrigeration
Entropy change accompanying specific
processes
Concentrating on the system:
1. The Helmholtz and Gibbs energies
2. Standard molar Gibbs energies.
Combining the First and Second Law:
1. Fundamental equation
2. Properties of internal energy
3. Properties of the Gibbs energy
The laws of thermodynamic describe the
basic properties of energy
 The first law of thermodynamics (also called law of
conservation of energy): the total amount of energy (E) within a
given system remains constant.
Energy can change form, such as from chemical E  heat E
 The second law of thermodynamics:
 E is converted from one form to another, the amount of useful E
decreases.
 All spontaneous changes result in a more uniform distribution of E,
reducing the E differences that are essential for doing work; E is
spontaneously converted from more useful into less useful forms.
A spontaneous reaction occurs naturally and favors the formation of products
at the specified conditions.
A spontaneous process is one that occurs without ongoing outside
intervention (such as the performance of work by some external force.
Chemical
potential
Direction of
spontaneous
change
A nonspontaneous reaction is a reaction that does not favor the formation
of products at the specified conditions.
Photosynthesis is a
nonspontaneous reaction
that requires an input of
energy.
The recognition of two classes of process, spontaneous and non-spontaneous is
summarized by the Second Law Thermodynamic.
This Law can be expressed in a variety of equivalent ways. One statement was
formulated by Kelvin:
“No process is possible in which the sole result is absorption of heat from a
reservoir and its complete conversion into work”
The Kelvin statement of the Second
Law denies that possibility of the
process illustrated here, in which heat
is changed completely into work,
there being no further change.
The process is not in conflict with the
First Law because energy is
conserved.
Example: Heats Engine
How to understand the role of distribution of energy?
Inelastic losses in the
materials of the ball
and floor.
Kinetic energy →energy
of the thermal motion
The direction of spontaneous change for a ball bouncing on a floor.
On each bounce some of its energy is degraded into the thermal motion of the
atoms of the floor, and that energy disperses.
The reverse has never been observed to take place on microscopic scale.
The molecular interpretation of the reversibility expressed by the Second Law.
(a) A ball resting on a warm surface
are undergoing thermal motion
(vibration, in this instance),
as indicated by the arrows
(b) For the ball to fly upwards, some of the
random vibrational motion would have to
change into coordinated, direction motion.
Such as conversion is highly improbable.
Entropy
A thermodynamic function that is proportional to the number of energetically
equivalent ways to arrange the components of a system to archive a particular
state; a measurement of the energy randomization or energy dispersal in a system.
Entropy symbol: S.
The First Law uses the internal energy to identify permissible changes.
The Second Law uses the entropy to identify the spontaneous changes among those
permissible changes.
The entropy of thermodynamic can be expressed in terms of the entropy:
The entropy of an isolated system increases in the course of a spontaneous change
∆Stot>0
Where Stot is the total entropy of the system and its surroundings.
Thermodynamically irreversible process (like cooling the temperature of the
surroundings and the free expansion of gases) are spontaneous processes, and hence
must be accompanied by increase in total enthalpy.
(a) The Thermodynamic Definition of Entropy
The thermodynamic of entropy concentrates on the change in entropy, dS that occurs as a
result of a physical or chemical change (in general, as a result of a ‘process’).
The thermodynamic definition of entripy is based on the expression:
Definition of
entropy change
Eqn 3.1
Where qrev is the heat supplied reversibly. For measureable change between two states
i and f this expression integrates to:
Eqn 3.2
That is, to calculate the difference in entropy between any two states of a system, we find
a reversible path between them, and integrate the energy supplied as heat at each stage
of the path divided by the temperature at which heating occurs.
Example:
Calculating the entropy change for the isothermal expansion of a perfect gas.
Calculate the entropy change of a sample of a perfect gas when its expand
isothermally from value Vi to a volume Vf.
Answer:
Because the temperature is constant, Eqn 3.2 becomes
From equation
We know that
It follows that
Therefore, we can adapt the definition in Eqn 3.1, delete the constraint “reversible”, and write
Entropy change of the
surroundings
Eqn 3.3
Furthermore, because the temperature of the surroundings is constant whatever the change,
for the measurable change
Eqn 3.4
That is, regardless of how the change is brought about in the system, reversibly or irreversibly
we can calculate the change of entropy of the surroundings by dividing the heat transferred
at which the transfer take place.
Eqn 3.3 makes it very simple to calculate the changes in entropy of the surroundings that
accompany any process. For instance, for any adiabatic change qsur=0, so
for adiabatic change:
∆Ssur = 0
This expression is true however the change takes place, reversibly or irreversibly, provided
No local hot spots are formed in the surroundings.
That is, it is so long as the surroundings remain in internal equilibrium. If hot spots do form,
then the localized energy may subsequently disperse spontaneous and hence generate
more entropy.
Example:
To calculate the entropy change in the surrounding when 1.00 mol H2O (l) is formed
from its elements under standard conditions at 298K, we use ∆Hɵ = -286kJ.The energy
released as is supplied to the surroundings, now regarded as being at constant
pressure, so qsur=+286kJ. Therefore,
This strongly exothermic reaction results in an increase in the entropy of the surroundings
as energy is released as heat into them.
(b) The Statistical View of Entropy
The entry point into the molecular interpretation of the Second Law of thermodynamic is
Boltzmann’s insight, that an atom or more molecule can possess only certain values of the
energy, called its “energy level”.
The continuous thermal agitation that molecules experience in a sample at T>0 ensures
that they are distributed over the available energy levels.
Boltzman made the link between the distribution of molecules over energy levels and the
Entropy. He propose that the entropy of a system is given by
Boltzmann formula for the entropy
Eqn 3.5
Where k=1.381 x 10-23 JK-1 and W is the number of microstate, the ways in which the
molecules of a system can be arranged while keeping the total energy constant.
Eqn 3.4 is known as Boltzman formula and the entropy calculated from it is sometimes called
the statistical entropy.
If W=1 (which corresponds to one microstate) then S=0 because ln 1=0. However if the
system can exist in more than one microstate, then W>1 and S>0.
(c) The Entropy As A State Function
Entropy is state function.
To prove this assertion, we need to show
that the integral of dS is independent to the path.
To do so, it is sufficient to prove that the integral
of Eqn 3.1 around an arbitrary cycle is zero,
for that guarantees that the entropy is the
same at the initial and final states of the system
regardless of the path taken between them.
That is, we need to show that
Eqn 3.6
In the thermodynamic cycle, the overall
change in a state function (from the initial
state to the final state and then back to the
initial state again) is zero.
Where the symbol denotes integration around a closed path. There are three steps in
The argument:
1. First, to show that Eqn 3.6 is true for a special cycle (a “Carnot Cycle”) involving
a perfect gas.
2. Then to show that the result is true whatever the working substance.
3. Finally, to show that the result is true for any cycle.
In Step 1, there is isothermal reversible
expansion at the temperature Th.
Step 2 is reversible adiabatic expansion in
which the temperature falls from Th to Tc.
Step 3 there is an isothermal reversible
compression at Tc, and that isothermal
reversible step is followed by an adiabatic
reversible compression,which restores the
system to its initial state.
The basic structure of a Carnot Cycle
The total change in entropy around the cycle is the sum of the changes in each of
these four step:
However, we show in the following Justification that for the perfect gas
Eqn 3.7
Substitution of this relation into the proceeding equation gives zero on the right,
which is we wanted to prove.
In second step we need to show Eqn 3.6 applies to any material, not just a perfect gas
(which is why, in anticipation, we have not labelled it with ao).
We begin this step by introducing the efficiency, ᶯ (eta) of a heat engine:
Definition of
efficiency
Eqn 3.8
Suppose an energy qh (for example, 20kJ)
is supplied to the engine and qc is lost from
the engine (for example,qc=-15kJ) and
discarded into the cold reservior.
The work done by the engine is equal to qh+qc
(for example, 20kJ+(-15kJ)=5kJ).
The efficiency is the work done divided by the
energy supplied as heat from the hot source.
Eqn 3.9
It then following from Eqn 3.7 (nothing that the modulus signs remove the minus sign)
that
Carnot efficiency
Eqn 3.10
(a) The demonstration of the equivalent of the efficiencies
of all reversible engines working between the same
thermal reservoirs is based on the flow of energy
represented in this diagram.
(b) The effect of the processes is the conversion of heat
into work without there is being a need for a cold sink;
this is contrary to the Kelvin statement of the Second Law.
The Second Law of thermodynamic
implies that all reversible engines have
the same efficiency regardless of their
construction.
A general cycle can be divided into small
Carnot cycles. The match in the limit of
infinestimally small cycle.
Path cancel in the interior of the collection,
and only the perimeter, an increasingly good
approximation to the true cycle as the number
of cycles increases, survives.
Because the entropy change around every
individual cycle is zero, the integral of the
entropy around the perimeter is zero too.
Therefore, all the entropy changes cancel except for those along parameter of the
overall cycle. That is,
The Thermodynamic Temperature
Suppose we have an engine that is working reversibly between a hot source at a
Temperature Th and a cold sink at a temperature T, then we know from Eqn 3.10 that
Eqn 3.11
The Clausius inequality
Clausius
inequality
Eqn 3.12
Refrigerators, Air conditioners
and Heat Pumps
When energy migrates from cool
source at a temperature Tc into a
warmer sink at a temperature Th,
the change in entropy is
Eqn 3.14
Eqn 3.15
Definition of coefficient of
performance
(d) Entropy Changes Accompanying Specific Processes
(a) Expansion
Entropy change for the
isothermal expansion of a perfect
gas
Eqn 3.17
Because S is a state function, the value of ∆S of the system is independent of
the path between the initial and final states, so this expansion applies whether
the change of state occurs reversibly or irreversibly.
The logarithmic increase in entropy of a
perfect gas it expands isothermally.
For any process the energy lost as heat from the system is acquired by the
surroundings, so dqsur=-dq. For a reversible change we use the expansion
In (qrev=nRT ln(Vf/Vi).
Eqn 3.18
This change is the negative of the change in the system, so we can conclude
that ∆Stot=0, which is what we should expect for a reversible process.
If, on the other hand, the isothermal expansion occurs freely (w=0), then q=0
(because ∆U=0). Consequently, ∆Ssur=0, and the total entropy change is given by:
In this case, ∆Stot>0, as we expect for an irreversible process.
(b) Phase transition
Consider a system and its surroundings at the normal transition temperature, Ttrs,
The temperature at which two phases are in equilibrium at 1 atm.
This temperature is 0oC (273K) for ice in equilibrium with liquid water at 1 atm, and
100oC (373K) for liquid water in equilibrium with its vapour at 1 atm.
At the transition temperature, any transfer of energy as heat between the system
and its surroundings is reversible because the two phase in the system are
equilibrium. Because at constant pressure q=∆trsH, the change in molar entropy of
system is:
Entropy of phase
transition
Eqn 3.20
List some experimental entropies of transition.
List in more details the standard entropies of
vaporization of several liquids at their boiling
points.
(c) Heating
We can use eqn 3.2 to calculate the entropy of a system at a temperature Tf from
a knowledge of its entropy at another temperature Ti and the heat supplied to charge
its temperature from one value to the other:
Eqn 3.21
Then from the definition of constant-pressure heat capacity
Entropy variation with
temperature
Eqn 3.22
The same expression applies at constant volume, but with Cp replaced by Cv.
When Cp is independent of temperature in the temperature range of interest, it can
be taken outside the integral and obtain
With a similar expression for heating at constant volume.
The logarithmic increase in entropy of a substance as it is heated at constant volume.
Difference curves correspond to different values of the constant-volume heat capacity
(which is assumed constant over the temperature range) expressed as Cv,m/R.
(d) The Measurement of Entropy
The entropy of a system at a temperature T is related to its entropy at T=0 by
measuring its heat capacity Cp at different temperatures and evaluating the integral
In Eqn 3.22, taking care to add the entropy of transition (∆trsH/Ttrs) for each phase
transition between T=0 and the temperature of interest.
Eqn 3.24
All the properties required, except Sm(0), can be measured calorimetrically, and the
integrals can be evaluated either graphically or, as is now more usual, by fitting
a polynomial to the data and integrating the polynomial analytical.
The variation of Cp/T with the
temperature for a sample is used to
evaluate the entropy, which is equal
to the area beneath the upper curve
up to the corresponding temperature,
plus the entropy of each phase
transition passed.
One problem with the determination of entropy is the difficulty of measuring heat
capacities neat T=0. There are good theoretical grounds for assuming that the
heat capacity is proportional to T3 when T is low and the dependence is the basis
of the Debye extrapolation.
In this method, Cp is measured down to as low a temperature as possible, and a curve
of the form aT3 is fitted to the data. That fit determines the value of a, and the
expression Cp,m=aT3 is assumed valid down to T=0.
The Helmholtz and Gibbs Energies
The Clausius inequality implies a number of criteria for spontaneous change under
a variety of conditions that may be expressed in terms of the properties of the system
alone; they are summarized by introducing the Helmholtz and Gibss energy.
A spontaneous process at constant temperature and volume is accompanied by a
decrease in the Helmholtz energy.
The change in the Helmholtz energy is equal to the maximum work accompanying
a process at constant pressure.
A spontaneous process at constant pressure and pressure is accompanied by the
decrease in the Gibbs energy.
The change in is equal to the maximum non-expansion work accompanying a process
at constant temperature and pressure.
Consider a system in thermal equilibrium with its surroundings at a temperature T.
When a change in a system occurs and there is a transfer of energy as heat between
the system and surroundings, the Clausius inequality (ds≥dq/T) reads
Eqn 3.27
We can develop this inequality in two ways according to the conditions (of constant
volume or constant pressure) under which the process occurs.
(a)Criteria for Spontaneous
First consider heating at constant volume. Then, in the absence of non-expansion
work, we can write dqv=dU, consequently
Eqn 3.28
The importance of inequality in this form is that it expresses the criterion for
spontaneous charge solely in terms of the state function of the system. The inequality
is easily rearrange into
Tds≥dU (constant V, no additional work)
Eqn 3.29
At either constant internal energy (du=0) or constant entropy (dS=0), this expression
Become, respectively,
dSU,V ≥0 dUS,V
Eqn 3.30
Where the subscirpts indicate the constant conditions.
When energy is transferred as heat or constant pressure, and there is no work other
than expansion work, we can write dqp=dH and obtain
TdS≥dH
(constant p,no additional work)
Eqn 3.31
At either constant enthalpy or constant entropy this inequality becomes, resprectively
dSH,p≥0 dHS,p≤0
Eqn 3.32
The interpretation of these inequality are similar to those of eqn 3.30.
The entropy of the system at the constant pressure must increase if its entalpy remains
constant (for they can then be no change in entropy of the surroundings).
Alternatively the enthalpy must decrease if the entropy of the system is constant, for
then is essential to have an increase in entropy of the surroundings.
Because Eqn 3.29 and 3.31 have the forms dU-TdS≤0 and dH-TdS≤0, respectively
they can expressed more simply by introducing two more thermodynamic quantities.
A=U-TS
G=H-TS
Definition of
Helmholtz Energy
Definition of Gibbs
energy
Eqn 3.31
Eqn 3.32
All the symbols in these two definition refer to the system.
When the state of the system changes at constant pressure, the two properties
change as follows:
(a) dA=dU-TdS
(b) dG=dH-TdS
Eqn 3.35
The criteria of the spontaneous change as
(a) dAT,V≤0
(b) dGT,p≤0
Eqn 3.36
These inequality are the most important conclusions from thermodynamic for chemistry.
(b) Some remarks on the Helmholtz energy
A change in a system at constant temperature and volume is spontaneous if dAT,V≤0.
The criterion of equilibrium when neither the forward nor reverse process has a
tendency to occur, is
dAT,V=0
Eqn 3.37
The expression dA=dU-TdS and dA=0 are sometimes interpreted as follows :
-ve dA,-ve dU, +ve TdS
This is observation suggests that the tendency of a system to move to a lower A is due
to its tendency to move towards states of lower internal energy and higher entropy.
However, this interpretation is FALSE the tendency to lower A is solely a tendency
towards states of greater overall entropy.
System change spontaneously if in doing so the total entropy of the system and its
surroundings increase, not because they tend to lower internal energy.
The form of dA may give the impression that the systems favoured lower energy, but
misleading: dS is the entropy change of the system, -dU/T is the entropy change of the
surroundings (when the volume of the system is constant), and their total tends to a
maximum.
(c) Maximum work
The change in the Helmholtz function is equal to the maximum work accompanying
a process at constant temperature:
dwmax=dA
Eqn 3.38
As a result, A is sometimes called the “maximum work function” or the “work function”
Justification of Maximum work:
To demonstrate that maximum work can be expressed in terms of the changes in
Helmholtz energy, we combine the Clausius inequality dS≥dq/T in the form TdS≥dq
with the first Law, dU=dq + dw and obtain
dU≤TdS + dw
dU is smaller than the term of the right because we are replacing dq by TdS which
in general is larger) this expression rearranges to
dw≥dU – TdS
It follows that the most negative value of dw and therefore the maximum energy
that can be obtained from the system as work, is given by
dwmax = dU – TdS
And that this work is done only when the path is traversed reversibly (because then
the equality applies). Because at constant temperature dA =dU-TdS, we conclude
that dwmax=dA.
with
Eqn 3.39
Eqn 3.40
Relation between A and
maximum work
In a system not isolated from its surroundings, the work done may be
different from the change in internal energy. Moreover, the process is
spontaneous if overall the entropy of the system and its surroundings
increase. In the process depicted here, the entropy of the system
decrease, so that of the process to be spontaneous, which means that
energy must pass from the system to the surrounding as heat. Therefore
less work than ∆U can be obtained.
In these process, the entropy of the system increase; hence we can
afford to lose some entropy of the surroundings. That is some of their
energy can be returned to them as work. Hence the work done can
exceed ∆U
(d) Some remarks on the Gibbs Energy
The criterion dGT,p≤0 carries over into chemistry as the observation that, at constant
temperature and pressure, chemical reactions are spontaneous in the direction
of decreasing Gibbs energy.
If G decreases as the reaction proceeds, then the reaction has a spontaneous
tendency to convert the reactant into products.
If G increases, then the reverse reaction is spontaneous.
In such reactions, H increases the system rises spontaneously to states of higher
enthalpy, and dH>0. Because the reaction is spontaneous we know that dG<0 despite
dH>0; it follows that the entropy of the system increase so much that TdS outweigh
dH in dG = dH –TdS.
Endothermic reactions are therefore driven by the increase of entropy of the system,
and this entropy change overcomes the reduction of entropy brought in the
surroundings by the inflow of heat into the system (dSsur= -dH/T at constant pressure).
(e) Maximum Non-Expansion Work
The analogue of the maximum work interpretation of ∆A and the origin of the name
“free energy” can be found for ∆G. In the following Justification, we show that at
Constant temperature and pressure the maximum additional (non-expansion) work,
Wadd,max is given by the change in Gibbs energy:
dwadd,max=dG
Eqn 3.14(a)
The corresponding ecpression for measureable change is
wadd,max =∆G
Relation between G and
maximum non-expansion work
Eqn 3.14(b)
Justification of Maximum non-expansion work:
Because H=U+pV, for general change in condition the change in enthalpy is
dH = dq + dw + d(pV)
The corresponding change in Gibbs energy (G=H-TS) is
dG = dH –TdS –SdT = dq + dw + d(pV) – TdS – SdT
When the charge is isothermal we can set dT = 0; then
dG = dq + dw + d(pV) –TdS
When the change is reversible, dw = dwrev and dq = dqrev = TdS, so for a reversible,
isothermal process
dG = TdS + dwrev + d(pV) – TdS = dwrev + d(pV) – TdS = dwrev + d(pV)
The work consists of expansion work, which is for a reversible change is given by
-pdV and possibly some other kind of work (for instance, the electrical work of pushing
electrons through a circuit of raising a column of liquid); this additional work we donate
dwadd. Therefore with d(pV) = pdV + Vdp,
dG= (-pdV + dwadd,rev) + pdV + Vdp =dwadd,rev +Vdp
If the change occurs at constant pressure (as well as constant temperature) we can set
dp=0 and obtain dG=dwadd,rev.
Therefore, at constant temperature and pressure, dwadd,rev = dG. However, because
the process is reversible, the work done must now have its maximum value so
Eqn 3.41 follows.
Standard Molar Gibbs Energy
Standard Gibbs energies of formation are used to calculate the standard Gibbs
energies of reactions. The Gibbs energies of formation of ions may be estimated from
a thermodynamic cycle and the Born equation.
Definition of standard Gibbs
energy of reaction
Eqn 3.42
Procedure for calculating
the standard Gibbs energy
of reaction
Eqn 3.43 (a)
Eqn 3.43 (b)
Convention for ions solution
Eqn 3.44
Born equation
Eqn 3.45a
Where Zi is the charge number of the ion and
ri its radius (NA is Avogadro’s constant).
Note that ∆solvGɵ <0 and that ∆solvGɵ is
strongly negative for small, highly charged
ions in media of high relative permittivity. For
water for which εr=78.54 at 25oC,
Eqn 3.45b
Combining the First and Second Law
The Fundamental Equation
The First Law of Thermodynamic may be written dU = dq + dw.
For a reversible change in a closed system of constant composition, and in the
absence of any additional (non-expansion) work, we may be set dwrev=-pdV and
(from the definition of entropy) dqrev=TdS, where p is the pressure of the system
and T its temperature. Therefore, for a reversible change in a closed system,
dU = TdS – pdV
The fundamental equation
Eqn 3.46
dU is exact differential, its value is independent of path.
Therefore, the same value of dU is obtained whether the change is brought about
irreversibly or reversibly.
Consequently, Eqn 3.46 applies to any change – irreversibly or reversibly – of a
closed system that does no additional (non-expansion) work.
We shall call this combination of the First and Second Law the Fundamental
Equation.
Properties of The Internal Energy
Relationship between thermodynamic properties are generated by combining
thermodynamic and mathematical expression for changes in their value.
The Maxwell relationship are a series of relations between derivatives of
thermodynamic properties based on criteria for changes in the properties being
exact differentials.
The Maxwell relationship are used to derive the thermodynamic equation of state
and to determine how the internal energy of a substance varies with volume.
The mathematical consequence of U being a function of S and V is that we can
express an infinitesimal change dU in terms of changes dS and dV by
Eqn 3.47
The two partial derivatives are the slope of the plots of U against S and V, respectively.
When this expression is compared to the thermodynamic relation, Eqn 3.46 we see
that for system of constant composition,
Eqn 3.48
(a) The Maxwell relations
An infinitesimal change in a function
can be written
where g and h are function of x and y. The mathematical criterion for df being an exact
differential (in the sense that its integral is independent of path) is that
Eqn 3.49
Because the fundamental equation, Eqn 3.46 is an expression for an exact differential,
the functions multiplying dS and dV (namely T and –p) must pass this test. Therefore,
It must be the case
A Maxwell
relation
Eqn 3.50
(b) The Variation of Interval Energy With Volume
The quantity
which represents how the internal energy charges
as the volume of a system is change isothermally, played a central role in the
manipulation of the First Law and we used the relation
A thermodynamic
equation of state
Eqn 3.51
This relation is called thermodynamic equation of state because it is an expression
for the pressure in terms of a variety of thermodynamic properties of the system.
We now ready to derive it by using a Maxwell relation.
Properties of The Gibbs Energy
(a) General consideration
When the system undergoes a change of state, G may change because H,T and S
all change.
dG = dH – d(TS) = dH – TdS – SdT
Because H=U+pV, we know that
dH = dU + d(pV) = dU + pdV + Vdp
And therefore,
dG = dU + pdV + Vdp - TdS – SdT
For a closed system doing no non-expansion work, we can replace dU by the
fundamental equation dU=TdS-pdV and obtain
dG = TdS – pdV + pdV + Vdp – TdS - SdT
Four terms now cancel on the right, and we conclude that for a closed system in
absence of non-expansion work and at constant composition
dG = Vdp – SdT
The fundamental equation of
chemical thermodynamic
The variation of G
with T and P
Eqn 3.52
Eqn 3.53
The vibration of the Gibbs energy of
a system with (a) temperature at constant
pressure and (b) pressure constant
temperature.
The slope of the formal is equal to the
negative of the entropy of the system and
that of the latter is equal to the volume.
The variation of the Gibbs energy with the temperature is determine by the entropy.
Because the entropy of the gaseous phase of a substance is greater than of that
liquid phase, and the entropy of the solid phase is smallest, the Gibbs energy
changes most steeply for the gas phase, followed by the liquid pahse, and then
the solid phase of the substance.
(b) The variation of The Gibbs Energy With Temperature
The first relationship in Eqn 3.53,
is our stating point for this
discussion. Although it expresses the variation of G in terms of the entropy, we can
express it in terms of the enthalpy by using the definition of G to write S=(H-G)/T.
Then
Eqn 3.54
We shall see later that the equilibrium constant of a reaction is related to G/T
rather than to G itself, and its it is easy to deduce from the last equation
Gibbs –Helmholtz
equation
Eqn 3.55
This expression is called the Gibbs-Helmholtz equation. It show that, if we have the
enthalpy of the system, then we know how G/T varies with temperature.
(c) The Variation of The Gibbs Energy With Pressure
To find the Gibbs energy at one pressure in terms of its value at another pressure,
the temperature being constant, we set dT=0 in Eqn 3.52 which gives dG=Vdp and
integrate:
Eqn 3.57a
For molar quantities,
Eqn 3.57b
This expression is applicable to any phase of matter, but to evaluate it we need to
know how molar volume, Vm, depends on the pressure.
The molar volume of a condensed phased changes only slightly as the pressure
changes, so we can treat Vm as a constant and take it outside the integral:
Eqn 3.58
The difference in Gibbs energy of a solid or liquid at two pressure is equal to the
rectangular area shown. We have assumed that the variation of volume with
pressure is negligible.
The molar value of gases are large, so Gibbs energy of a gas depends strongly on the
pressure. Furthermore, because the volume also varies markedly with the pressure, we
cannot treat in the integral in eqn 3.57b.
For a perfect gas we substitute Vm=RT/p into the integral, treat RT as constant and find:
Eqn 3.59
This expression shows that when the pressure is increase tenfold at room temperature,
the molar Gibbs energy increase by RT ln 10≈9 kJmol-1. It also follows from this
equation that, if we set pi=pɵ (the standard pressure of 1 bar), then the molar Gibbs
energy of a perfect gas at a pressure p(set pf=p) is related to its standard value by
The molar Gibbs energy
of a perfect
Eqn3.60
The difference in Gibbs energy for
a perfect gas at two pressures is equal
to the area shown below the perfectgas isotherm.
The molar Gibbs energy of a perfect
gas is proportional to ln p, and the
standard state is reached at pɵ. Note
that, as p→ 0, the molar Gibbs
energy becomes negatively infinite.
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